1. Energy Changes in Reactions

2. Learning objectives  Perform simple energy calculations using different energy units  Apply specific heat concept to heat calculations  Distinguish between endothermic and exothermic reactions  Calculate specific heat from calorimetry  Calculate heat changes in calorimetry

3. Energy is capacity to do work  Energy comes in different forms  Kinetic energy is energy due to motion  Potential energy is energy due to position or state

4. Energy comes in many forms  Kinetic energy is energy due to motion  Heat  Rotation  Vibration  Translation  Potential energy is energy due to position or state  Height  Chemical  Electrical  Coiled spring

5. Energy is interchangeable  Processes convert energy from one form to another  Chemical reaction (potential → heat)(potential → heat)  Turbine (heat → mechanical)(heat → mechanical)  Generator (mechanical → electrical)(mechanical → electrical)  Light bulb (electrical → light)(electrical → light)  Photosynthesis (light → chemical)(light → chemical)  Some energy is always wasted during conversion

6. Measuring energy: calories, Calories and joules  calorie is energy required to raise temperature of 1 g of water 1 degree C  Calorie is the food version = 1,000 cal  Raises temperature of 1 pint of water 3.8ºF  Joule is SI unit derived from mechanical work: work done when force of 1 Newton is applied for 1 meter 1 cal = 4.18 J

7. Specific heat  Specific heat is energy required to raise temperature of 1 g of substance by 1ºC  Units are J/gºC or cal/gºC

8. Champagne from a styrofoam cup: Measuring heat change  Calorimetry is process used to measure heat change of reaction  Heat change in calorimeter is: Q = m H2O x ΔT H2O x SH H2O Weigh Measure Know

9. Principle of Calorimetry  Conservation of energy:  Exothermic:  Temperature in calorimeter increases  Heat lost by process (system) = heat gained by H 2 O (surroundings)  Endothermic  Temperature in calorimeter decreases  Heat gained by process = heat lost by H 2 O

10. Exo-thermic and endo-thermic  H 2 + O 2 gives out energy – exothermic  The system (chemical bonds) lose potential energy)  N 2 + O 2 absorbs energy – endothermic  The system gains potential energy

11. Calorimetry example: Calculating specific heat  What is specific heat of lead if 57.0 J are required to raise temperature of 35.6 g Pb by 12.5ºC? Q = m s x ΔT s x SH s

12. Specific heat metal example  35.2 g metal at 100ºC are placed in calorimeter containing 42.5 g H 2 O at 19.2ºC. Final temperature is 29.5ºC. What is SH of metal? SH H 2 O = 4.18 J/g ºC  Heat lost by metal = heat gained by H 2 O  Q metal = mass metal x ΔT metal x SH metal  Q water = mass water x ΔT water x SH water

13. Heat of reaction  Heat of reaction is the enthalpy change when reactants are converted into products  Note sign:  Exothermic (energy out) is –  Endothermic (energy in) is +

14. Nutrients and energy  Macronutrients supply energy  Fats (9 Cal/g) (Main form of energy storage)  Carbohydrates (4 Cal/g)  Proteins (4 Cal/g)